Chem 1721 Brief Notes: Chapter 19

Transcription

1 Chem 1721 Brief Notes: Chapter 19 Chapter 19: Electrochemistry Consider the same redox reaction set up 2 different ways: Cu metal in a solution of AgNO 3 Cu Cu salt bridge electrically conducting wire Ag Ag+ NO3 - Cu2+ Ag+ Anode Cathode What is the reaction? In the experiment on the left: over time the strip of copper metal decreases in size (and mass) the solution starts clear and colorless, and over time becomes clear and light blue over time a fine grayish-silver powder deposits on the bottom of the beaker ox: Cu (s) Cu 2+ (aq) + 2 e red: Ag + (aq) + 1 e Ag (s) net: Cu + 2 Ag + Cu Ag energy change? Any energy associated with the process goes into changing the temperature of the solution. In the experiment on the right: the Cu/Cu 2+ redox couple and the Ag + /Ag redox couple are separated Cu metal and Ag metal are connected by an electrically conducting wire e transfer from Cu to Ag + occurs through the wire generating a voltage this is an electrochemical cell there are two types of electochemical cells 1. Galvanic cell: a spontaneous chemical reaction occurs that generates a voltage 2. Eletrolytic cell: a nonspontaneous chemical reaction is driven by an applied current Galvanic cells 2 compartments: anode and cathode anode where oxidation occurs; usually shown on the left cathode where reduction occurs; usually shown on the right can write anode and cathode half reactions (parallel to oxidation and reduction half reactions respectively) anode: cathode: net cell: Cu Cu e Ag e Ag Cu + 2 Ag + Cu Ag each compartment must contain an electrically conducting solid; the electrode; where the wire connects in the example above: copper is the anode, silver is the cathode the direction of electron flow in a Galvanic cell is always anode cathode electrons are generated at the anode (where oxidation occurs) electrons are consumed at the cathode (where reduction occurs)

4 the half reaction that is lower in the table (larger E ) will be the cathode (reduction) half reaction the half reaction that is higher in the table (smaller E ) will be the anode (oxidation) half reaction remember: must end up with a + E cell for a Galvanic cell relative strengths of oxidizing and reducing agents the lower in the table the stronger the oxidizing agent can cause any half reaction above it to proceed in reverse F 2 is the strongest oxidizing agent listed Li is the strongest reducing agent listed a related concept... the Activity Series of Metals ranks metals as reducing agents metals that like to be oxidized are good reducing agents good reducing agents have small (and frequently negative) E s Activity Series: Li > K > Ba > Ca > Mg > Be > Al > Zn > Fe > Cu > Ag > Au most active metal least active metal best reducing agent worst reducing agent on the list a metal in the Activity Series can reduce any M n+ to the right of it lots of potential questions several examples Consider these half reactions: NO H e NO + 2 H 2 O E = 0.96 V Fe 3+ + e Fe 2+ E = 0.77 V If combined in a Galvanic cell, what will be the net cell reaction and E cell? Can MnO 4 in acidic solution (i.e H + present) oxidize Ni? Ag? Which is a stronger reducing agent, Cr or Mn? Based on the Acitivity Series, can aluminum metal reduce Ca 2+? Ag +? Complete description of a Galvanic cell at this point you should be able to write/identify the anode and cathode half reaction of a cell write the net cell reaction determine E cell, or E anode or E cathode if E cell is given sketch a diagram for a Galvanic cell including: identification of the anode and cathode compartments and their components (i.e electrode and ions in solution); direction of electron flow; salt bridge or porous barrier and ion migration write the line notation that describes a Galvanic cell

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